Which of the following best describes the nature of metallic bonding?

The nature of metallic bonding is best described by overlapping atomic orbitals leading to the delocalization of electrons. In metallic bonds, metal atoms lose some of their electrons, which become delocalized and can move freely throughout the metal lattice. This delocalization creates a "sea of electrons" that allows for the conductivity of electricity and heat, as well as other characteristic metallic properties such as malleability and ductility.

The delocalized electrons are not associated with any specific atom but instead contribute to the overall metallic structure, which explains the cohesive nature and strength of metallic bonds. The presence of these mobile electrons allows metals to exhibit properties such as luster and the ability to deform without breaking, as the electrons can reorient themselves as the metal undergoes stress.

In contrast, the other options do not accurately capture the characteristics of metallic bonding. Atoms arranged in rigid geometrical forms and discrete independent molecules refer more to covalent or ionic character in network solids or molecular solids, while atoms with no orbital interactions would suggest a lack of bonding altogether, which is not applicable in the context of metallic substances. Thus, delocalization of electrons is the defining feature that characterizes metallic bonding effectively.